theory and practice
A catalyst is a
substance that alters the rate of chemical reaction without itself
being permanently chemically changed.
It will chemically change
temporarily e.g. change in ligand or oxidation state or other bonding
arrangement, but will return to is original state often via a 2–3
stage 'catalytic cycle'.
A catalyst provides a
reaction pathway with a lower activation energy Ea,
compared to the uncatalysed reaction (see diagrams immediately below, simple
exothermic/endothermic reaction and more realistically, a complex two
stage cycle profile further on).
Two primary modes
of catalytic action – heterogeneous and homogeneous
HETEROGENEOUS CATALYSIS: (e.g. diagram above
for nickel catalysing the hydrogenation of an alkene)
The catalyst and
reactants are in different phases (usually solid catalyst and
The reaction occurs
on the catalyst surface which may be the transition metal or one
of its compounds e.g. an oxide.
The reactants must be
adsorbed onto the catalyst surface at the 'active sites'.
This can be physical
adsorption or chemically bonding to the catalyst surface. Either
way, it has the effect of concentrating the reactants close to
each other and weakening the original intra–molecular bonds of
the reactant molecules.
The diagram above
illustrates a typical heterogeneous catalysis e.g. hydrogenation of
alkenes with hydrogen and a nickel catalyst.
The strength of
adsorption is crucial to having a 'fruitful' catalyst surface.
adsorption too strong,
the reactant/product molecules are too strongly 'chemisorped'
inhibiting reaction progress e.g. can happen with tungsten (W).
adsorption too weak, the
reactants are not chemisorped strongly enough to allow the
initial bond breaking processes to happen e.g. can happen with silver (Ag),
though silver is used in some industrial processes.
Nickel (Ni), platinum (Pt), rhodium (Rh) etc. will adsorb the
reactants sufficiently to enable the bond breaking process
to be initiated but to not strong to retain the product
molecules. These three metals are used in many industrial
processes e.g. hydrogenating oils to make margarine (Ni) and
catalytic converters in vehicle exhausts (Pt, Rh).
It is usual
the catalyst in a finely divided form to maximise surface area to
give the greatest and therefore most efficient rate of reaction.
This means the catalyst must be physically supported. since it
will have no bulk strength in its own right e.g.
should be avoided. This inhibiting effect is caused by
impurity molecules being strongly chemisorbed on the most active sites
of the catalyst surface. It considerably reduces the efficiency of
the catalyst and increases production costs if the catalyst has to
be replaced or functions with less efficiency e.g.
the iron catalyst in the Haber Process for making ammonia,
and lead poisons the
platinum–rhodium surface in car exhaust catalytic converters.
The catalyst and
reactants are in the same phase (usually a solution), and so the
catalysed reaction can happen throughout the bulk of the reaction
is usually due to temporary changes in oxidation state of a
transition metal ion and results in a 'catalytic
cycle'. In other words, the homogeneous catalysed reactions occur via some
e.g. (i) Either iron(II) Fe2+ ions or iron(III)
Fe3+ ions catalyse the oxidation of iodide ions by
in diagram above) the overall
(aq) + 2I– (aq) ==> 2SO42–
(aq) + I2 (aq)
'direct' uncatalysed reaction involves the collision of two
repelling negative ions and so has a high activation energy.
Activation energies arise from outer electron shell
repulsions and bond energies.
collision of an Fe3+ ion and an I– ion
involves positive–negative attraction which helps overcome the
repulsion component activation energy due to two negative
for catalysed (Ea1 in diagram above)
followed by (Ea2
in diagram above)
The iron(III) ion
is regenerated in the cycle whether you start with Fe2+
or Fe3+ , showing the iron ions act in a
genuine catalytic way and the iron ions are not consumed
overall in the process.
you added up the two equations of the cycle you get the
equation of overall reaction change.
is an excellent example of why transition element
compounds can act as catalysts in specific redox
reactions i.e. they can exist in, and interchange
between, two (or more) oxidation states that facilitate
the overall reaction.
Note 3: Eø arguments can be used to check on the
feasibility of the reaction or mechanism steps.
autocatalysis by Mn2+ ions when the oxidising
agent potassium manganate(VII), KMnO4, is used to
titrate the ethanedioate ion, C2O42–,
(from acid/salt, old name 'oxalic/oxalate').
+ 16H+(aq) + 5C2O42–(aq)
==> 2Mn2+(aq) + 8H2O(l) +
the reactant collisions are between two anions which will
have a high activation energy, hence the slow start to the
reaction. In the titration you see it gradually gets
faster and faster because there is catalytic cycle
involving the hexa–aqua Mn(II) ion and ethanedioate
complexes of Mn(II) and Mn(III).
[MnII(H2O)6]2+(aq) ==> [MnII(C2O4)3]4–(aq) ==>
[MnIII(C2O4)3]3–(aq) ==> [Mn(H2O)6]2+(aq) + CO2(aq/g)
that the catalytic cycle involves changes in ligand and
oxidation state in the manganese metal ions and two
ions catalyse the oxidation of the 2,3–dihydroxybutandioate
ion (acid/salt, old name 'tartaric/tartrate') to water, methanoate
ion and carbon dioxide with hydrogen peroxide solution.
The likely scheme of events is outlined below, the
equations are NOT
meant to be balanced.
pink hexa–aqa Co2+ ion, which is a Co(II)
pink Co(II) complex changes ligand
from water to the organic acid, but no change in oxidation
state or co–ordination number, and I
don't know its colour?, but it perhaps it doesn't
exist long enough to be seen?
[Co(OOCCH(OH)CH(OH)COO)3]3–(aq) ==> [Co(H2O)6]2+(aq),H2O(l),HCOO–(aq),CO2 (aq/g)
green Co(III) complex then breaks down to
give the products,
you see the bubbles of carbon dioxide and the 'return'
pink hexa–aqa Co2+ complex ion.
the above sequence, the change in ligand affects the
relative stability of the oxidation states. The CoII–acid
complex is stable as regards 'breakdown', but is readily
oxidised to the CoIII–acid complex, which is
NOT stable to breakdown.
* Titanium * Vanadium
* Manganese * Iron * Cobalt
* Copper *
* Silver & Platinum
Introduction 3d–block Transition Metals * Appendix
Hydrated salts, acidity of
hexa–aqua ions * Appendix 2. Complexes
& ligands * Appendix 3. Complexes and isomerism * Appendix 4.
Electron configuration & colour theory *
Appendix 5. Redox
equations, feasibility, Eø * Appendix 6.
Catalysis * Appendix 7.
* Appendix 8. Stability Constants and entropy
Appendix 9. Colorimetric analysis
and complex ion formula * Appendix 10 3d block
– extended data
* Appendix 11 Some 3d–block compounds, complexes, oxidation states
& electrode potentials * Appendix 12
Hydroxide complex precipitate 'pictures',
formulae and equations